Sodium Catches Fire in Cold Water While Copper Ignores Steam

Drop a small piece of sodium into cold water and it fizzes across the surface and catches fire. Drop a piece of copper into boiling water and nothing happens at all — and nothing happens in steam either.

Both are metals. Both have loose outer electrons. The difference is how readily each one gives those electrons away, and that single quantity — reactivity — decides everything in this part of the chapter:

- whether a metal reacts with oxygen, water or acid, and how violently
- whether one metal can displace another from its salt
- which extraction method is needed to obtain it from its ore

So the chemistry of metals is essentially one ordered list applied over and over. That list is the reactivity series, and by the end of this page you should be able to predict a reaction you have never seen simply by looking up two metals' positions in it.

The chapter then asks what happens when a metal meets a non-metal: the metal loses electrons, the non-metal gains them, and the two oppositely charged ions hold together in an ionic bond. Almost every property of common salt follows from that one transfer.

This page covers the second part of the CBSE Class 10 Science chapter on metals and non-metals: reactions with oxygen, water and acids, the reactivity series, ionic bond formation, and the properties of ionic compounds.

Metal with oxygen gives an oxide; metal with water gives a hydroxide or an oxide plus hydrogen; metal with a dilute acid gives a salt plus hydrogen.

With oxygen. All these are combination reactions:

$$4\text{Na}+{\text{O}}_2\to 2{\text{Na}}_2\text{O}$$
$$2\text{Mg}+{\text{O}}_2\to 2\text{MgO}$$
$$4\text{Al}+3{\text{O}}_2\to 2{\text{Al}}_2{\text{O}}_3$$
$$2\text{Cu}+{\text{O}}_2\to 2\text{CuO}$$

Sodium and potassium react so fast at room temperature that they are stored under kerosene; magnesium needs to be lit; copper needs strong heating; and silver and gold do not react with oxygen at all.

Metal oxides are basic, and the soluble ones give alkalis:

$${\text{Na}}_2\text{O}+{\text{H}}_2\text{O}\to 2\text{NaOH}$$

Two oxides react with both acids and bases, and these are called amphoteric:

$${\text{Al}}_2{\text{O}}_3+6\text{HCl}\to 2{\text{AlCl}}_3+3{\text{H}}_2\text{O}$$
$${\text{Al}}_2{\text{O}}_3+2\text{NaOH}\to 2{\text{NaAlO}}_2+{\text{H}}_2\text{O}$$

Zinc oxide behaves the same way. Amphoteric oxides are a favourite one-mark question, and aluminium and zinc are the two names to remember.

With water. The violence of the reaction falls steadily down the series:

$$2\text{Na}+2{\text{H}}_2\text{O}\to 2\text{NaOH}+{\text{H}}_2\text{(cold water, so exothermic the hydrogen ignites)}$$
$$\text{Ca}+2{\text{H}}_2\text{O}\to {\text{Ca(OH)}}_2+{\text{H}}_2\text{(cold water, but calmly)}$$
$$\text{Mg}+2{\text{H}}_2\text{O}\to {\text{Mg(OH)}}_2+{\text{H}}_2\text{(hot water only)}$$
$$3\text{Fe}+4{\text{H}}_2\text{O}\to {\text{Fe}}_3{\text{O}}_4+4{\text{H}}_2\text{(steam only)}$$

and copper, silver and gold react with none of them.

A detail worth knowing about calcium. It reacts more slowly than sodium, yet the piece of calcium starts to float — because bubbles of hydrogen stick to its surface and lift it. The floating is a clue, not the reaction, and questions ask for the reason.

With dilute acids. Metal plus dilute acid gives a salt and hydrogen:

$$\text{Mg}+2\text{HCl}\to {\text{MgCl}}_2+{\text{H}}_2$$
$$\text{Zn}+{\text{H}}_2{\text{SO}}_4\to {\text{ZnSO}}_4+{\text{H}}_2$$
$$\text{Fe}+2\text{HCl}\to {\text{FeCl}}_2+{\text{H}}_2$$

and the rate falls in the order magnesium, zinc, iron — while copper, silver and gold give no hydrogen at all, because they are below hydrogen in the series.

One acid behaves differently and it is examined. Nitric acid is a strong oxidising agent, so instead of releasing hydrogen it usually oxidises the metal and produces oxides of nitrogen. Magnesium and manganese with very dilute nitric acid are the only common exceptions that do release hydrogen. So a question asking which acid does not give hydrogen with copper has the answer none of them, but a question asking why does dilute nitric acid not give hydrogen with zinc wants the oxidising-agent reason.

A metal higher in the series displaces one lower down from its salt solution. If it is lower, nothing happens.

The series, from most to least reactive:

$$\text{K}>\text{Na}>\text{Ca}>\text{Mg}>\text{Al}>\text{Zn}>\text{Fe}>\text{Pb}>(\text{H})>\text{Cu}>\text{Hg}>\text{Ag}>\text{Au}$$

Hydrogen is placed in it deliberately: a metal above hydrogen releases hydrogen from a dilute acid, and a metal below it does not.

Worked example 1. Will iron displace copper from copper sulphate solution?

Iron is above copper, so yes:

$$\text{Fe(s)}+{\text{CuSO}}_4\text{(aq)}\to {\text{FeSO}}_4\text{(aq)}+\text{Cu(s)}$$

The blue solution fades towards pale green and a reddish-brown deposit appears on the nail — two visible changes that confirm the prediction.

Worked example 2. Will copper displace iron from iron sulphate solution?

Copper is below iron, so no reaction occurs. The solution stays as it was, and no reaction is the complete answer.

Worked example 3. Will silver displace copper from copper sulphate?

Silver is below copper, so no. Displacement runs one way only — down the series, never up.

Worked example 4 — with a stronger metal. Zinc in copper sulphate:

$$\text{Zn(s)}+{\text{CuSO}}_4\text{(aq)}\to {\text{ZnSO}}_4\text{(aq)}+\text{Cu(s)}$$

Worked example 5 — displacement from an oxide. Aluminium is above iron, so it can pull the oxygen away from iron oxide:

$${\text{Fe}}_2{\text{O}}_3\text{(s)}+2\text{Al(s)}\to 2\text{Fe(l)}+{\text{Al}}_2{\text{O}}_3\text{(s)}$$

This releases so much heat that the iron comes out molten, which is how cracked railway tracks and machine parts are welded on site. The reaction is a displacement and a redox at the same time — aluminium is oxidised and iron oxide reduced.

Why the series predicts so much from so little. A metal high in the series loses its electrons easily. That is why it reacts fast with oxygen, water and acid, why it can push a less willing metal out of a compound, and — as Part 3 shows — why it is hardest to extract from its ore in the first place. One ordering, four consequences.

The boundary case that tests understanding. Potassium and sodium are so far above hydrogen that they react with water itself, so they cannot be used for displacement in an aqueous solution — they attack the water before reaching the salt. A reaction can be too vigorous to be useful, and that is why copper is displaced with zinc or iron rather than with sodium.

One atom hands over electrons and the other accepts them, and the two resulting ions are held together by electrostatic attraction.

Sodium chloride. Sodium has the electron arrangement $2,8,1$ and chlorine has $2,8,7$.

- Sodium loses its single outer electron and becomes ${\text{Na}}^{+}$ with the arrangement $2,8$ — the same as neon
- Chlorine gains that electron and becomes ${\text{Cl}}^{-}$ with the arrangement $2,8,8$ — the same as argon
- The oppositely charged ions attract, giving $\text{NaCl}$

$$\text{Na}\to {\text{Na}}^{+}+e^{-}\text{Cl}+e^{-}\to {\text{Cl}}^{-}$$

Both ions now have a completely filled outer shell, which is why the transfer happens at all: losing one electron is easier for sodium than gaining seven, and gaining one is easier for chlorine than losing seven.

Magnesium chloride. Magnesium has the arrangement $2,8,2$, so it must lose two electrons:

- Magnesium loses two electrons to become ${\text{Mg}}^{2+}$ with the arrangement $2,8$
- Two chlorine atoms each gain one of them, becoming two ${\text{Cl}}^{-}$ ions
- The formula is therefore ${\text{MgCl}}_2$

The formula came out of the electron count, not from memory. Whenever you can say how many electrons each atom gains or loses, the formula follows, and that is the quickest way to check a formula you are unsure of.

Worked check — formula unit masses. With the atomic masses Na $=23$, Mg $=24$ and Cl $=35.5$:

$$\text{NaCl}=23+35.5=58.5{\text{MgCl}}_2=24+2(35.5)=95$$

**Note the words formula unit mass rather than molecular mass.** An ionic compound has no molecules at all: a crystal of common salt is a vast three-dimensional lattice of alternating sodium and chloride ions, with no single $\text{NaCl}$ unit that can be picked out. The formula tells you the ratio, not a particle — and writing a molecule of sodium chloride is a standard error.

Why only metals and non-metals do this. A metal has few outer electrons and loses them readily; a non-metal is a few electrons short of a full shell and accepts them readily. The transfer needs one of each, which is why two metals never bond ionically and why two non-metals share electrons instead — the covalent bonding of the next chapter.

Because the ions are locked in a rigid lattice by strong electrostatic forces, and they can carry current only once those forces let them move.

Every property of an ionic compound follows from that one sentence.

High melting and boiling points. The electrostatic attraction between oppositely charged ions is strong, and every ion is attracted by several neighbours at once. Melting means breaking a large number of those attractions, which needs a great deal of energy — so common salt melts only at a very high temperature, while a covalent substance such as wax melts in a warm pan.

Hard and brittle. The lattice resists being squeezed, so the crystals are hard. But push the layers sideways and ions of like charge are forced next to each other; they repel, and the crystal splits cleanly along a plane. Hardness and brittleness together are the signature of an ionic solid — a metal, by contrast, bends because its electron sea lets the layers slide.

Solubility. Ionic compounds generally dissolve in water and not in petrol or kerosene. Water molecules surround each ion and pull it away from the lattice; a non-polar solvent cannot. That is why salt dissolves in water and grease does not, and why a spilt oil stain cannot be washed out with water alone.

Electrical conductivity, which is the most examined property.

- As a solid, an ionic compound does not conduct. The ions are charged, but they are fixed in the lattice and cannot move
- Molten, it conducts. Heat has broken the lattice, so the ions are free to drift towards the electrodes
- Dissolved in water, it conducts. The lattice has been pulled apart by the water, so again the ions are free

**So the requirement is not charged particles but mobile charged particles. Solid salt has the charges and lacks the mobility; molten salt has both. That distinction is the whole answer** to why does solid sodium chloride not conduct electricity while its solution does, and an answer that only says because it has ions earns nothing.

Worked comparison. Put two electrodes into dry salt crystals and the bulb stays dark. Add water and it lights. Dry the same salt out again and the bulb goes dark once more. Nothing chemical has changed — only whether the ions can move.

One boundary case. A few ionic compounds are only slightly soluble, and some covalent compounds ionise in water and do conduct — hydrochloric acid is the obvious example from the last chapter. Conductivity in solution proves mobile ions, not an ionic bond, and the two questions are different.